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In inorganic chemistry, nitrites are salts of nitrous acid (HNO2). They contain the nitrite ion (NO2−). Nitrites of the alkali and alkaline earth metals can be synthesized by reacting a mixture of nitrogen monoxide (NO) and nitrogen dioxide (NO2) with the corresponding metal hydroxide solution, as well as through the thermal decomposition of the corresponding nitrate. Other nitrites are available through the reduction of the corresponding nitrates.
Sodium nitrite is used for the curing of meat because it prevents bacterial growth and, in a reaction with the meat's myoglobin, gives the product a desirable dark red color. Because of the toxicity of nitrite (the lethal dose of nitrite for humans is about 22 mg per kg body weight), the maximum allowed nitrite concentration in meat products is 200 ppm. Under certain conditions, especially during cooking, nitrites in meat can react with degradation products of amino acids, forming nitrosamines, which are known carcinogens.
Nitrite is detected and analyzed by the Griess Reaction, involving the formation of a deep red-colored azo dye upon treatment of a NO2−-containing sample with sulfanilic acid and naphthyl-1-amine in the presence of acid.
In organic chemistry, nitrites are esters of nitrous acid and contain the nitrosooxy functional group. They possess the general formula RONO, where R is an aryl or alkyl group. Amyl nitrite is used in medicine for the treatment of heart diseases.
Nitrites should not be confused with nitrates, the salts of nitric acid, or with nitro compounds, though they share the formula RNO2. The nitrite anion NO2− should not be confused with the nitronium cation NO2+.
- V. M. Ivanov (2004). "The 125th Anniversary of the Griess Reagent". Journal of Analytical Chemistry. 59 (10): 1002–1005. Translated from V. M. Ivanov (2004). Zhurnal Analiticheskoi Khimii. 59 (10): 1109–1112. Missing or empty